Glossary: Chemical Bonding Key Terms

Term	Definition
Ionic bond	Electrostatic attraction between oppositely charged ions formed by electron transfer from metal to non-metal
Covalent bond	Bond formed by sharing of electron pair(s) between two atoms
Sigma (σ) bond	Bond formed by head-on (axial) overlap of orbitals along the internuclear axis; one per single/double/triple bond
Pi (π) bond	Bond formed by lateral (sideways) overlap of parallel p orbitals; present in double (1π) and triple (2π) bonds
Hybridization	Mixing of atomic orbitals of similar energies to form a new set of equivalent hybrid orbitals
Steric number (SN)	Number of σ bonds + lone pairs on the central atom; determines hybridization
VSEPR	Valence Shell Electron Pair Repulsion theory; predicts geometry by minimising electron pair repulsions
Dipole moment (μ)	Product of charge (q) and distance (d): μ = q×d; unit: Debye (D); measures polarity
Lattice energy (U)	Energy required to separate one mole of ionic solid into gaseous ions; always endothermic
Born-Haber cycle	Thermochemical cycle applying Hess's Law to indirectly calculate lattice energy
Fajan's rules	Rules predicting covalent character: increases with small cation, large anion, high cation charge
MOT	Molecular Orbital Theory; uses LCAO to form bonding (σ,π) and antibonding (σ,π) MOs
Bond order (BO)	$\frac{Nb−Na}{2}$ ; measures net bonding; 0 = molecule doesn't exist; higher BO = shorter, stronger bond
Paramagnetic	Substance with unpaired electrons; weakly attracted to magnetic field
Diamagnetic	Substance with all electrons paired; weakly repelled by magnetic field
Resonance	Delocalization of electrons across multiple equivalent Lewis structures; actual molecule is the hybrid
Resonance energy	Extra stability of resonance hybrid compared to most stable single resonance structure
Hydrogen bond	Intermolecular/intramolecular electrostatic attraction: X–H···Y (X,Y = F, O, or N)
Metallic bond	Electrostatic attraction between positive metal ions and delocalized electron sea
Madelung constant	Geometrical factor in lattice energy formula; depends on crystal structure type
Electron affinity (EA)	Energy change when a gaseous atom/ion gains an electron; $EA_{1}$ is usually exothermic for non-metals
Ionisation energy (IE)	Energy required to remove an electron from a gaseous atom/ion; always endothermic
Polar molecule	Molecule with μ ≠ 0; has a net dipole due to asymmetric charge distribution
s-p mixing	Interaction between σ2s and σ2p MOs in elements with Z ≤ 7; raises energy of σ2p above π2p